CHE202-02 S25 Chapter 8 notes

Course Information

• Course Title: Chemical Bonding Course: General Chemistry and Qualitative Analysis II

• Course Code: CHE202-02

• Schedule: Tuesday and Thursday: 9:30 – 10:45 am

• Instructor: Dr. Kesete Ghebreyessus or Dr. Ghebre

• Textbook: Chemistry the Central Science, 15th Edition, by Brown, LeMay, and others

Chapter 8: Concepts of Chemical Bonding

• Definition of Chemical Bond:

  • The force that holds different atoms together in a molecule.

8.1 Types of Chemical Bonds

• Three Basic Types of Bonds:

  • Ionic Bonds:

    • Attraction between charged ions.

  • Covalent Bonds:

    • Sharing of electrons between atoms.

  • Metallic Bonds:

    • Metal atoms bonded to several other atoms.

8.2 Ionic Bonding

• Formation of Ionic Bonds:

  • Involves the transfer of one or more electrons from one atom to another.

  • Typically occurs between metals and nonmetals.

  • Example Reaction:

    • Na(s) + 1/2 Cl2(g) → NaCl(s)

  • Electron Configurations:

    • Na: [Ne]3s1 (Valence electron)

    • Cl: [Ar]3s23p5

Energetics of Ionic Bond Formation

• Heat of Formation:

  • The reaction to form sodium chloride is exothermic.

  • DHf° = -411 kJ/mol

• Lattice Energy:

  • Energy required to completely separate a mole of solid ionic compound into its gaseous ions.

  • Formula: Eel = κ (Q1Q2/d)

    • Q1 and Q2: charges of the ions.

    • d: distance between the ions, related to the radius/size.

  • Endothermic Reaction:

    • Separation of NaCl into sodium and chloride ions: DH = +788 kJ/mol.

Lattice Energy Factors

• Increases with:

  • Higher charges (Q1 and Q2) on the ions.

  • Smaller ionic size.

Example of Lattice Energy

• Arrange following ionic compounds in order of increasing lattice energy:

  • a) NaCl, KCl, LiCl, CsCl

  • b) LiF, CsBr, MgO, CaO

8.3 Covalent Bonding

• Definition of Covalent Bonding:

  • Formed by sharing electrons between atoms.

• Lewis Dot Symbols:

  • Represents chemical symbols and valence electrons.

• Electron Configuration Example (Oxygen):

  • O: [He]2s22p4

• Chlorine Electron Configuration:

  • Cl: [Ne]3s23p5 (seven valence electrons).

Octet Rule

• Atoms react to achieve a full octet (eight valence electrons).

• Lewis Structures:

  • Illustrate the formation of covalent bonds, showing shared and unshared electron pairs.

Multiple Bonds

• More than one pair of electrons can be shared (e.g., O2 forms a double bond).

• Bond Length and Strength:

  • Triple bonds are shorter and stronger than double bonds.

  • Double bonds are shorter and stronger than single bonds.

Bond Polarity and Electronegativity

• Electronegativity (EN):

  • Ability of an atom to attract electrons.

  • EN increases left to right across periods, decreases top to bottom in groups.

• Polar Covalent Bonds:

  • Bonds with unequal electron sharing due to different electronegativities.

  • EN Difference:

    • DEN ≥ 0.5 and < 2.0 = polar.

    • DEN < 0.5 = non-polar.

    • DEN > 2.0 = ionic compound.

8.5 Drawing Lewis Structures

1. Sum valence electrons of all atoms (add one for each negative charge, subtract one for each positive charge).

2. Identify the central atom and connect outer atoms with single bonds.

3. Complete octets of outer atoms, tracking electrons remaining.

4. Complete octet of the central atom.

5. If octet is not complete, create multiple bonds.

8.6 Formal Charge (FC)

• FC Calculation:

  • FC = Valence Electrons - (Lone Pair Electrons + 1/2 Bonding Electrons).

• The total FC should equal the overall charge of the molecule.

8.7 Resonance Structures

• Resonance:

  • Different Lewis structures for the same molecule indicating delocalized electrons (e.g., O3, SO3, benzene).

Exceptions to the Octet Rule

1. Molecules with fewer than eight electrons (e.g., BF3).

2. Molecules with more than eight electrons (expanded octet - e.g., PF5).

3. Molecules with an odd number of electrons (e.g., NO).

Summary of Key Concepts

• Importance of ionic and covalent bonding in chemistry.

• Application of the octet rule and Lewis structures for predicting molecular geometry and stability.

• Understanding of electronegativity and its role in bond polarity.